`=>` Werner was the first to describe the bonding features in coordination compounds.

`=>` But his theory could not explain following properties of coordination compounds :

(i) remarkable property of forming coordination compounds by certain elements

(ii) directional properties of bonds in coordination compounds

(iii) characteristic magnetic and optical properties of coordination compounds

`=>` Many approaches have been put forth to explain the nature of bonding in coordination compounds viz. Valence Bond Theory (VBT), Crystal Field Theory (CFT), Ligand Field Theory (LFT) and Molecular Orbital Theory (MOT). We shall discuss only application of VBT and CFT to coordination compounds.

`=>` According to this theory, the metal atom or ion under the influence of ligands can use its `color{red}((n-1)d, ns, np)` or `color{red}(ns, np, nd)` orbitals for hybridisation to yield a set of equivalent orbitals of definite geometry such as octahedral, tetrahedral, square planar and so on (Table 9.2).

`=>` These hybridised orbitals are allowed to overlap with ligand orbitals that can donate electron pairs for bonding. This is illustrated by the following examples.

`=>` It is usually possible to predict the geometry of a complex from the knowledge of its magnetic behaviour on the basis of the valence bond theory.

`color{red}("Example ")` : (i) In the diamagnetic octahedral complex, `color{red}([Co(NH_3)_6]^(3+))`, the cobalt ion is in `+3` oxidation state and has the electronic configuration `3d^6`. The hybridisation scheme is as shown in diagram.

● Six pairs of electrons, one from each `color{red}(NH_3)` molecule, occupy the six hybrid orbitals. Thus, the complex has octahedral geometry and is diamagnetic because of the absence of unpaired electron.

● In the formation of this complex, since the inner `color{red}(d)`-orbital (`color{red}(3d)`) is used in hybridisation, the complex, `color{red}([Co(NH_3)_6]^(3+))` is called an `text(inner orbital)` or `color{green}(("low spin")` or `color{green}("spin paired complex")`.

(ii) The paramagnetic octahedral complex, `color{red}([CoF_6]^(3–))` uses outer orbital (`color{red}(4d)`) in hybridisation `color{red}(sp^3 d^2)`.

● It is thus called `color{green}("outer orbital")` or `color{green}("high spin")` or `color{green}("spin free complex")`.

(iii) In tetrahedral complexes one `color{red}(s)` and three `color{red}(p)` orbitals are hybridised to form four equivalent orbitals oriented tetrahedrally.

● This is illustrated below for `color{red}([NiCl_4]^(2-))`. Here nickel is in `+2` oxidation state and the ion has the electronic configuration `3d^8`.

● The hybridisation scheme is as shown in diagram.

● Each `color{red}(Cl^-)` ion donates a pair of electrons.

● The compound is paramagnetic since it contains two unpaired electrons.

● Similarly, `color{red}([Ni(CO)_4])` has tetrahedral geometry but is diamagnetic since nickel is in zero oxidation state and contains no unpaired electron.

(iv) In the square planar complexes, the hybridisation involved is `color{red}(dsp^2)`.

● Example : `color{red}([Ni(CN)_4]^(2–))`.

● Here nickel is in `+2` oxidation state and has the electronic configuration `color{red}(3d^8)`.

● The hybridisation scheme is as shown in diagram.

● Each of the hybridised orbitals receives a pair of electrons from a cyanide ion.

● The compound is diamagnetic as evident from the absence of unpaired electron.

`color{red}("Note ")` : It is important to note that the hybrid orbitals do not actually exist. In fact, hybridisation is a mathematical manipulation of wave equation for the atomic orbitals involved.

`=>` The magnetic moment of coordination compounds can be measured by the magnetic susceptibility experiments.

`=>` The results can be used to obtain information about the structures adopted by metal complexes.

`=>` For metal ions with upto three electrons in the `color{red}(d)`-orbitals, like `color{red}(Ti^(3+), (d^1); V^(3+) (d^2); Cr^(3+), (d^3))`; two vacant `color{red}(d)`-orbitals are available for octahedral hybridisation with `color{red}(4s)` and `color{red}(4p)` orbitals.

● The magnetic behaviour of these free ions and their coordination entities is similar.

`=>` When more than three `color{red}(3d)` electrons are present, the required pair of `color{red}(3d)` orbitals for octahedral hybridisation is not directly available (as a consequence of Hund’s rule).

● Thus, for `color{red}(d^4 , (Cr^(2+), Mn^(3+)), d^5 (Mn^(2+), Fe^(3+)), d^6 (Fe^(2+), Co^(3+)))` cases, a vacant pair of `color{red}(d)` orbitals results only by pairing of `color{red}(3d)` electrons which leaves two, one and zero unpaired electrons, respectively.

`=>` The magnetic data agree with maximum spin pairing in many cases, especially with coordination compounds containing `color{red}(d^6)` ions.

`=>` However, with species containing `color{red}(d^4)` and `color{red}(d^5)` ions there are complications.

`=>` `color{red}([Mn(CN)_6]^(3-))` has magnetic moment of two unpaired electrons while `color{red}([MnCl_6]^(3-))` has a paramagnetic moment of four unpaired electrons.

● `color{red}([Fe(CN)_6]^(3-))` has magnetic moment of a single unpaired electron while `color{red}([FeF_6]^(3-))` has a paramagnetic moment of five unpaired electrons.

● `color{red}([CoF_6]^(3-))` is paramagnetic with four unpaired electrons while `color{red}([Co(C_2O_4)_3]^(3-))` is diamagnetic.

`=>` This apparent anomaly is explained by valence bond theory in terms of formation of inner orbital and outer orbital coordination entities.

● `color{red}([Mn(CN)_6]^(3-), [Fe(CN)_6]^(3-))` and `color{red}([Co(C_2O_4)_3]^(3-))` are inner orbital complexes involving `color{red}(d^2sp^3)` hybridisation, the former two complexes are paramagnetic and the latter diamagnetic.

● On the other hand, `color{red}([MnCl_6]^(3-), [FeF_6]^(3-))` and `color{red}([CoF_6-]^(3-))` are outer orbital complexes involving `color{red}(sp^3d^2)` hybridisation and are paramagnetic corresponding to four, five and four unpaired electrons.

While the VB theory, to a larger extent, explains the formation, structures and magnetic behaviour of coordination compounds, it suffers from the following shortcomings :

(i) It involves a number of assumptions.

(ii) It does not give quantitative interpretation of magnetic data.

(iii) It does not explain the colour exhibited by coordination compounds.

(iv) It does not give a quantitative interpretation of the thermodynamic or kinetic stabilities of coordination compounds.

(v) It does not make exact predictions regarding the tetrahedral and square planar structures of 4-coordinate complexes.

(vi) It does not distinguish between weak and strong ligands.